Redox Titrations and Limitations of the Oxidation Number Concept

You have already seen how acid-base titrations use a pH-sensitive indicator to pinpoint the moment an acid has exactly neutralised a base. Redox chemistry has its own version of the same idea. In a redox titration, you use a solution of known concentration (either an oxidant or a reductant) to find the concentration of an unknown reductant or oxidant, and a redox-sensitive indicator marks the endpoint (the point at which the indicator signals that the reaction is complete). The logic is identical: keep adding the titrant (the standard solution added from the burette) until the reaction is exactly complete, note the volume used, and calculate from the stoichiometry (the quantitative mole ratio between reactants and products).

What makes redox titrations fascinating is the variety of clever ways chemists have devised to spot the endpoint. Some reagents advertise the endpoint all by themselves through their own colour, others rely on a separate indicator molecule, and one elegant method uses a chain of reactions involving iodine and starch. Let us look at each approach.

Self-Indicating Reagents: The Permanganate Example

Some oxidising agents are so intensely coloured that they serve as their own indicator, no external dye needed. The best-known example is the permanganate ion ($MnO_4^-$), found in potassium permanganate ($KMnO_4$). Its deep pink-purple colour is visible even at extremely low concentrations.

Here is how a permanganate titration works in practice:

• You place the reductant (say $Fe^{2+}$ or $C_2O_4^{2-}$) in the flask and fill the burette with a standard $KMnO_4$ solution.
• As you add permanganate drop by drop, it immediately reacts with the reductant and gets decolourised. The solution in the flask stays colourless (or very faintly coloured) because every $MnO_4^-$ ion is consumed as fast as it arrives.
• The moment the last trace of reductant is used up, the very next drop of permanganate has nothing to react with. It remains in solution, giving the flask a faint but lasting pink tinge.

That first persistent pink colour is the endpoint. It appears at a $MnO_4^-$ concentration as low as $10^{-6}$ mol $L^{-1}$, which means there is almost no overshoot past the equivalence point (the point where the oxidant and reductant have reacted in exact stoichiometric proportions). This extreme sensitivity is what makes permanganate such a popular titrant.

Because no separate indicator is added, $MnO_4^-$ is called a self-indicator.

External Indicators: The Dichromate and Diphenylamine Pair

Not every oxidising agent has a colour dramatic enough to flag the endpoint on its own. Potassium dichromate ($K_2Cr_2O_7$) is a powerful oxidant, but its orange colour does not vanish or appear sharply enough at the equivalence point to serve as a reliable self-indicator.

The solution is to add a small amount of an external redox-sensitive indicator, a substance whose own oxidation produces a vivid colour change. For dichromate titrations, diphenylamine (an organic compound used as a redox indicator) is the classic choice:

• While reductant is still present in the flask, the dichromate reacts with it and leaves diphenylamine untouched.
• The instant the last trace of reductant is consumed, the very next bit of excess dichromate oxidises the diphenylamine instead.
• This oxidation produces an intense blue colour, which appears suddenly and unmistakably, signalling the endpoint.

The principle is straightforward: the indicator is chosen so that it reacts with the titrant only after all of the analyte (the substance whose concentration is being determined) has been used up. Until then, the analyte “protects” the indicator by consuming the titrant first.

The Iodometric Method: A Chain of Clever Reactions

The third approach, called iodometry (a method of analysis based on the liberation and measurement of iodine), is the most ingenious. It works for any reagent that is capable of oxidising iodide ions ($I^-$) to molecular iodine ($I_2$). Instead of looking for a colour change directly in the main reaction, chemists set up a chain of reactions that produces a sharp and unmistakable signal.

Step 1: The Main Reaction Liberates Iodine

Consider $Cu^{2+}$ as the oxidant. When copper(II) ions are mixed with excess iodide ions, the following redox reaction occurs:

$2Cu^{2+}(aq) + 4I^-(aq) \rightarrow Cu_2I_2(s) + I_2(aq) \qquad \text{(7.59)}$

Copper(II) is reduced to cuprous iodide ($Cu_2I_2$, an insoluble white solid), and iodide is oxidised to molecular iodine ($I_2$). The amount of $I_2$ released is directly proportional to the amount of $Cu^{2+}$ that was present.

Step 2: Starch Signals the Presence of Iodine

Once iodine is liberated, a small amount of starch solution is added. Iodine reacts with starch to produce an intense blue colour. This blue colour is unmistakable and highly sensitive, confirming that free iodine is present in the solution.

A practical note: although $I_2$ is insoluble in pure water, it dissolves readily in the presence of potassium iodide ($KI$). The iodine combines with $I^-$ to form the soluble tri-iodide ion, which effectively stays in solution as $KI_3$.

Step 3: Thiosulphate Removes the Iodine

Now a standard solution of sodium thiosulphate ($Na_2S_2O_3$) is added from the burette. Thiosulphate reacts specifically with iodine in another redox reaction:

$I_2(aq) + 2S_2O_3^{2-}(aq) \rightarrow 2I^-(aq) + S_4O_6^{2-}(aq) \qquad \text{(7.60)}$

In this reaction, iodine ($I_2$) is reduced back to colourless iodide ($I^-$), and thiosulphate ($S_2O_3^{2-}$) is oxidised to tetrathionate ($S_4O_6^{2-}$, the four-sulphur product formed when two thiosulphate units join together).

Step 4: The Endpoint

As long as free iodine remains, the solution stays blue. Each drop of thiosulphate removes a little more iodine. The moment the last trace of iodine has been consumed by thiosulphate, the blue colour vanishes. That disappearance of the blue colour is the endpoint.

From the volume of thiosulphate used, you calculate the moles of $I_2$ that were present, and from that you work backwards through the stoichiometry to find the original amount of $Cu^{2+}$ (or whichever oxidant released the iodine). The arithmetic is the only step left.

Why Iodometry Is So Versatile

The beauty of this method is that it is not limited to copper(II). Any oxidant that can convert $I^-$ to $I_2$ can be analysed this way. The starch-iodine-thiosulphate detection system stays the same regardless of which oxidant produced the iodine. This makes iodometry one of the most widely used techniques in analytical chemistry.

Looking Beyond Oxidation Numbers: An Evolving Concept

Throughout this chapter, we have used oxidation numbers as the central tool for understanding redox reactions: identifying which species is oxidised, which is reduced, balancing equations, and classifying reaction types. Oxidation numbers are powerful and practical, but it is worth understanding their limitations.

The oxidation number approach treats electron transfer as if it were complete: one atom “has” the electrons, another “does not.” In reality, most chemical bonds are covalent, meaning electrons are shared rather than fully transferred. The sharing is often unequal (polar bonds), but it is still sharing. Assigning a whole-number oxidation state to an atom in a covalent molecule is therefore an approximation, a convenient bookkeeping device rather than an exact physical description.

Recognising this limitation, modern chemistry increasingly describes redox processes in terms of electron density:

• Oxidation is a decrease in electron density around the atom(s) involved.
• Reduction is an increase in electron density around the atom(s) involved.

This way of thinking captures what is really happening in a molecule more accurately. When a carbon atom bonds to a more electronegative oxygen atom, the electron cloud shifts toward oxygen, so carbon’s electron density decreases (it is effectively oxidised) even though no electron has been physically “removed” in the way the oxidation number model implies.

The concept of redox is still evolving. What began as a simple observation about oxygen addition and removal has grown through the electron-transfer picture, through oxidation numbers, and now into the electron-density framework. Each layer builds on the one before, and each gives us a more refined understanding of how atoms exchange and share their electrons during chemical reactions.